📘Chemistry
Short Questions
- a) What is 1st ionization energy? Give an example.
- b) Explain why sulfur has a lower first ionization energy than phosphorus.
- c) Why the elements in Group 13 to 17 are called p-block elements?
- d) What are the factors that affect electronegativity?
- e) What factors are responsible for the increasing reactivity of alkali metals as we move down the group?
- f) Why some of the elements show variable oxidation numbers while others do not?
- g) Identify the element which is in period 5 and group 15?
- h) Why oxides of sodium and magnesium are more ionic than the oxides of nitrogen and phosphorous?
- i) Give reason for the different chemical reactivities of Na and Mg toward oxygen and chlorine.
- j) Why the ionization energy of lithium is much lower than that of helium despite the fact that the nuclear charge of lithium is +3 and that of helium is +2?
- k) The ionization energy of Be (atomic no. 4) is higher than that of B (atomic no. 5), despite the fact that the nuclear charge of Be is +4 and that of B is +5.
- l) What is common in Na , Mg²⁺, Al³⁺, Ne⁰ and F⁻? Arrange them in increasing order of sizes.
- m) Consider the chlorides of sodium, magnesium, and phosphorus(V): NaCl, MgCl₂, and PCl₅ (i) Classify each of these chlorides as acidic, basic, or neutral. (ii) For each chloride, briefly explain the reason for your classification, referring to their behavior when dissolved in water.
- a) There are three orientations of p-orbital due to three values of magnetic quantum number. Justify it.
- b) 'I3' of Mg is much bigger than its 'I2'. Justify.
- c) Among the elements Li, K, Ca, S and Kr which one has the lowest first ionization energy? Which has the highest first ionization.
- d) Consider the electronic configuration of the potassium atom (atomic number 19). (i) Write the full electronic configuration of potassium using the s, p, d, f notation. (ii) Explain why the 4s subshell is filled before the 3d subshell in potassium, even though the principal quantum number of the 3d subshell is lower.
- e) i) An atom of element X has an atomic number of 17 and a mass number of 35. Determine the number of protons, neutrons, and electrons in this atom. (ii) If this element forms an ion with a charge of -1, how many protons, neutrons, and electrons will be present in the ion?
- f) In the ground state of mercury soHg: i. How many electrons occupy atomic orbitals with n = 3? ii. How many electrons occupy 4d atomic orbitals? iii. How many electrons occupy 4p2 atomic orbital? iv. How many electrons in the valence shell have spin “up” (s = -½)?
- g) The successive ionization energies for an unknown element are I1 = 896 kJ/mol, I2 = 1752 kJ/mol I3 = 14,807 kJ/mol I4 = 17,948 kJ/mol To which family in the periodic table, does the unknown element most likely belong?
- b) Consider the following ionization energies for aluminium: (i) Account for the trend in the values of the ionization energies. (ii) Explain the large increase from I3 to I4. (iii) List the four aluminium ions given in order of increasing size, and explain your ordering.
- h) (i) State the general order of filling orbitals up to the 4p subshell. (ii) Explain why the 4s subshell is filled before the 3d subshell, according to the Aufbau principle.
- i) Draw the orbital box diagram for the valence electrons of a phosphorus atom (atomic number 15), ensuring that your diagram adheres to Hund's rule and the Pauli Exclusion Principle.
- Define the following: (i) Dipole (ii) Bond order (iii) Permanent dipole- permanent dipole force (iv) London dispersion force
- Draw the Lewis (electron dot) structures for the following species: (i) HCN (ii) NCl (iii) CO (iv) O₃ (v) NO₂
- Draw the orbital structure of the CO₂ molecule in terms of VBT.
- Draw the Lewis structures and tell whether the following ions involve expanded octets. i) ClO₄⁻ ii) ICl₄⁻ iii) NH₄⁺ iv) I₃⁻
- The bond between K and Cl is ionic but that between Si and Cl is polar covalent. Explain why.
- SO₂ is a polar molecule but SO₃ not. Justify.
- Which of O₂²⁺ and O₂²⁻ would be paramagnetic? Give reason in the light of MOT.
- Which of the following bonds would be most polar? i) C-Cl ii) Si-F iii) Se-F
- Compare the bond energies of single, double, and triple bonds between the same two atoms (e.g., H-H, O=O, N=N). Explain the trend in terms of the number of shared electrons.
- Predict the shapes of sulfate ([SO₄]²⁻), borate ([BH₄]⁻) and tri-iodide ions ([I₃]⁻) according to the VSEPR model.
- Sketch the molecular orbital pictures of π(2p) and π*(2p).
- Sketch the hybrid orbitals and bond formation in PCl₃, SiCl₄, and NH₄⁺.
- Explain the difference between the formation of σ and π bonds in terms of VBT.
- How is the concept of mole derived from Avogadro’s number?
- Define the following terms with one example in each case. (a) Molar mass (b) Molar volume (c) Molar concentration
- What do you mean by molar volume of a gas? How Avogadro’s number is related with molar volume?
- 39 g of potassium and 56 g of iron have equal number of atoms in them. Justify.
- 4g of He, 17 g of NH₃ and 64 g of SO₂ occupy separately the volumes of 22.414 dm³ although the sizes and molecular masses of molecules of the three gases are very different from each other. Explain.
- Do you think that 1 mole of H₂ and 1 mole of NH₃ at 0 °C and 1 atm will have Avogadro’s number of particles?
- What is stoichiometry? Give the basic assumptions of stoichiometric calculations.
- What is a limiting reactant? How does it control the quantity of the product formed?
- Differentiate theoretical and actual yields. How is the percentage yield of a reaction calculated?
- What are the factors which are mostly responsible for the low yield of the products in chemical reactions?
- Explain, at the molecular level, why evaporation leads to a cooling effect.
- Explain why liquids with stronger intermolecular forces tend to have lower rates of evaporation at a given temperature compared to liquids with weaker intermolecular forces.
- One feels sense of cooling under the fan after bath. Explain.
- How a dynamic equilibrium is established during evaporation of a liquid in a closed vessel at constant temperature?
- The boiling point of water is different at Lahore and Murree hills.
- Discuss two significant consequences of the lower density of ice compared to liquid water in natural environments.
- Why B.P of a liquid increases when the external pressure rises?
- Mention four items in which liquid crystals are used.
- How do you differentiate between crystalline solids and amorphous solids?
- Propanone (CH3COCH3), propanol (CH3CH2CH2OH) and butane (CH3CH2CH2CH3) have very similar relative molecular masses. List them in the expected order of increasing boiling points. Explain your answer.
- Discuss how hydrogen bonding is responsible for the relatively high surface tension of water.
- What type of intermolecular forces will dominate in the following liquids? (i) NH3(ii) Ar(iii) CH3COCH3 (iv) CH3OH
- a. Differentiate between exothermic and endothermic reactions.
- b. What do you understand by the enthalpy of a system?
- c. Differentiate clearly between entropy (S) and Gibbs free energy (G).
- d. Distinguish clearly between standard enthalpy of reaction and standard enthalpy of formation.
- e. Define the following enthalpies and give one example of each. (i) Standard enthalpy of solution (ii) Standard enthalpy of hydration (iii) Standard enthalpy of atomization (iv) Standard enthalpy of combustion
- f. Explain why the lattice enthalpy of an ionic compound is typically a large negative value.
- g. What factors influence the magnitude of the lattice enthalpy?
- h. Explain why the enthalpy of hydration is always an exothermic process for gaseous ions. What are the main interactions responsible for the release of energy during hydration?
- i. For the reaction CH₄(g) + 2O₂(g) → CO₂(g) + 2H₂O(g) identify all the bonds that need to be broken and all the bonds that need to be formed to carry out a bond energy calculation of ΔH.
- j. For a reaction to be spontaneous, what is the required sign of the Gibbs free energy change (ΔG)? Under what conditions of enthalpy change (ΔH) and entropy change (ΔS) will a reaction always be spontaneous?
- k. The enthalpy of solution can be either positive or negative. Explain what a positive ΔH_soln and a negative ΔH_soln indicate about the energy changes during the dissolution process.
- l. Consider two ions with similar charges but different sizes, or similar sizes but different charges. Explain how the concept of charge density can be used to predict which ion will have a more exothermic enthalpy of hydration and why.
- a) What do you understand by the rate of a reaction?
- b) Give the difference between enthalpy change of reaction and energy of activation of reaction.
- c) Differentiate clearly between order and molecularity of a reaction.
- d) Why the instantaneous rate changes during a reaction?
- e) Briefly summarize the effects of temperature and surface area on the rates of reactions.
- f) Justify that the radioactive decay is always a first order reaction.
- g) A reaction is second order with respect to a reactant. How is the rate of reaction affected if the concentration is doubled and reduced to half?
- h) What is meant by half-life and what is it used for?
- i) Why does wood burn more rapidly in pure oxygen than in air?
- j) A catalyst lowers the activation energy of a chemical reaction. Illustrate it.
- k) The rate constant for a certain reaction is 3.5 × 10⁻⁴ s⁻¹ at 25°C. What is the order of the reaction? Explain based on the units of the rate constant.
- l) If the initial concentration of the reactant is 0.50 mol dm⁻³, calculate the initial rate of the reaction.
- How would the rate of this reaction change if the concentration of the reactant were doubled?
- A certain first-order reaction has a rate constant of 2.5 × 10⁻³ s⁻¹. Calculate the half-life of the reaction in minutes.
- A radioactive isotope decays by a first-order process with a half-life of 12 hours. Calculate the rate constant for the decay in s⁻¹.
- a. What is meant by the state of chemical equilibrium?
- b. Define reversible reaction. Give an example.
- c. The change of volume disturbs the equilibrium position for some of the gas phase reactions but not the equilibrium constant.
- d. Mention the characteristics of chemical equilibrium.
- e. Reversible reaction attains the position of equilibrium which is dynamic in nature and not static. Explain it.
- f. Why do the rates of forward reactions slow down when a reversible reaction approaches the equilibrium stage?
- g. Why ice at 0 °C can be melted by applying pressure without supply of heat from outside?
- h. Write two conditions of equilibrium constant.
- i. The reversible reaction: 2SO₂(g) + O₂(g) → 2SO₃(g) has come to equilibrium in a vessel of specific volume at a given temperature. Before the reaction began, the concentrations of the reactants were 0.060 mol/dm³ of SO₂ and 0.050 mol/dm³ of O₂. After equilibrium was reached, the concentration of SO₃ was 0.040 mol/dm³. What is the equilibrium concentration of O₂?
- a. Define the following with an example for each: i) Ionization constant ii) Solubility product iii) Common ion effect iv) Acid-base Indicator
- b. Differentiate between: i) Hydrolysis and dissolution ii) Acidic and basic buffer solutions
- c. Explain the concept of conjugate acid-base pairs. How are they related in terms of proton transfer?
- d. What is the relationship between the strength of an acid and the strength of its conjugate base?
- e. For the following three reactions, identify the reactants that are Arrhenius bases, Bronsted-Lowry bases, and/or Lewis bases. State which type(s) of bases each reactant is. Explain your answers. i) NaOH(s) ⇌ Na⁺(aq) + OH⁻(aq) ii) HF(aq) + H₂O(l) ⇌ F⁻(aq) + H₃O⁺(aq) iii) H⁺(aq) + NH₃(aq) ⇌ NH₄⁺(aq)
- f. An amphoteric substance can behave as either an acid or a base. Identify whether water behaves as an acid or a base in each of the following reactions. i) H₂O + HCl ⇌ H₃O⁺ + Cl⁻ ii) NH₃ + H₂O ⇌ NH₄⁺ + OH⁻ iii) HNO₃ + H₂O ⇌ H₃O⁺ + NO₃⁻ iv) CH₃COOH + H₂O ⇌ CH₃COO⁻ + H₃O⁺
- g. Which salt would you expect to dissolve more readily in acidic solution: Barium carbonate or Copper sulfide? Explain. (Ksp(BaCO₃) = 1.1 × 10⁻¹⁰, Ksp(CuS) = 6 × 10⁻³⁶)
- h. Why does common ion effect decrease solubility of a less soluble salt?
- i. State the basic principle of solubility product. Mention factors affecting solubility product.
- j. Prove by equations what happens when Na₂CrO₄ is added to saturated solution of PbCrO₄.
- k. According to the Lewis acid-base concept, boron trifluoride (BF₃) can act as an acid. Is this statement correct?
- l. If the concentration of hydrogen ions in a solution is 1 × 10¹⁵ M, what is the pH of the solution?
- a. How and why electrical double layer is formed?
- b. Why electrode potential of Cu is called reduction potential?
- c. What are the advantages of salt bridge in a galvanic cell?
- d. How can we predict the feasibility of a chemical reaction using the cell voltage?
- e. During electrolysis of aqueous NaCl, why Na is not liberated at the cathode?
- f. Calculate the Ox. No. of chromium (Cr) in the following compounds: (i) CrCl₃ (ii) Cr₂(SO₄)₃ (iii) Cr₂O₇²⁻
- g. The order of decreasing reactivity of metals based on their position is K > Mg > Zn > Fe > Cu. Write balanced chemical equations for the reactions that would occur (if any) when: i) Copper is added to a solution of magnesium sulfate. ii) Iron is added to a dilute solution of hydrochloric acid.
- h. Explain why some metals higher in the activity series can displace hydrogen from acids, while others lower in the series cannot.
- i. Calculate the number of Faradays required to deposit 108 g of Ag⁺; 63.5 g of Cu²⁺ and 27 g of Al³⁺
- j. In an electrolysis experiment, a current of 0.500 A was passed through a solution of AgNO₃ for 30.0 minutes. The mass of silver deposited on the cathode was found to be 0.503 g. Given that the molar mass of silver is 107.87 g mol⁻¹ and the charge on a silver ion is +1. Calculate the value of Avogadro's number (N_A) from this data.
- k. A cell is set up with a standard nickel electrode (Ni²⁺ + 2e⁻ → Ni E⁰ = -0.25 V) and a standard cobalt electrode (Co²⁺ + 2e⁻ → Co E⁰ = -0.28 V). i) Identify which metal will be the anode and which will be the cathode. Justify your answer. ii) Write the balanced overall cell reaction. iii) Calculate the standard cell potential (E⁰_cell).
- a. Define the following: i) Cycloalkanes ii) Isomerism iii) Conjugated dienes iv) Inductive effect
- b. Differentiate between: i) Aliphatic and Aromatic hydrocarbons ii) Homolytic and Heterolytic Fission iii) Electrophile and Nucleophile
- c. Explain why alkanes do not undergo addition reactions.
- d. How elimination reaction is considered the opposite of an addition reaction?
- e. Compare the carbocation stability in propene and 2-Butene.
- f. Given the molecular formula C₄H₈, list all possible structural isomers that are alkenes.
- g. When propene (C₃H₆) undergoes electrophilic addition with HBr, it forms 2-bromopropane as the major product. Explain why 2-bromopropane is favored over 1-bromopropane, using the concept of carbocation stability.
- h. Explain why conjugated alkenes may show different reactivity compared to isolated alkenes.
- i. Explain how inductive effects from alkyl groups stabilize carbocations in alkenes.
- j. Write the equation for each reaction. i. CH₃CH₂CH=CH₂ with H₂ (Ni catalyst) ii. CH₃CH=CH₂ with Cl₂ iii. CH₃CH₂CH=CHCH₂CH₃ with H₂O (H₂SO₄ catalyst)
- k. Write structural formulas for each of the following compounds: i) Isobutylene ii) 2,3,4,4-Tetramethyl-2-pentene iii) 2,5-Heptadiene iv) 4,5-Dimethyl-2-hexene v) Vinylacetylene vi) 1,3-Pentadiene vii) 1-Butyne viii) 3-n-Propyl-1,4-pentadiene ix) Vinyl bromide x) But-1-en-3-yne xi) 4-Methyl-2-pentyne xii) Isopentane
- 1. Write down names of the following compounds according to IUPAC system: i) H₃C–CH–CH(CH₂)₂CH₃ ii) CH₂–CH₂–CH₂–C=CH₃ iii) CH₂–CH₂–CH₂–C=CH₃ iv) CH₂–CH–CH–CH₃ v) CH₂–C–CH₂CH₂CH₃ vi) (CH₃CH₂)₂CH₃ vii) CH₂CH₂C(CH₂)₂CH(CH₂CH₂CH₃) viii) CH₂–CH–C=C–CH=CH₃ ix) (CH₃)C–CH₂–C(CH₂)₃ x) CH₂C(CH₂)₂(CH₂)CH₃ xi) H₃C–CH₂–CH–CH₂–CH₂CH₃ xii) (CH₃)₂CH–CH–CH(CH₂)₃ xiii) CH₂–CH–CH–CH₂–CH₂–CH₃ xiv) CH₂–CH–CH–CH–CH₂–CH₃
- a) List two reasons for the inertness of N₂.
- b) How is nitrogen isolated from air?
- c) Why ammonia (NH₃) is a weak base?
- d) Write down the reactions of photochemical smog formation.
- e) What is the construction and function of a catalytic converter?
- f) Why sulfur is quite unreactive at room temperature?
- g) Which are the most stable oxidation states of sulfur in water at pH=0 and pH=14?
- h) How does sulfur react with halogens?
- i) Draw the structures of cyclo-octasulfur (S₈) and sulfuric acid.
- j) What is the role of sulfur in the vulcanization of rubber?
- k) What is the composition and the chemical reaction of gunpowder combustion?
- l) What is the importance of disulfide bridges?
- m) Write down self-ionization equation of sulfuric acid and its ionization in water.
- n) Give two examples where sulfuric acid acts as a dehydrating agent.
- o) How does V₂O₅ catalyze the formation of SO₃?
- p) What is purpose of formation of oleum?
- a. Which halogen is the least reactive, which is the most reactive? Give reason.
- b. The ionic equation for a reaction is: Br₂ + 2I⁻ → 2Br⁻ + I₂ Explain which species is oxidized in this reaction. Why?
- c. What is role of London dispersion forces in the trend of volatility of halogens?
- d. How does the reactivity of halogens with hydrogen vary?
- e. Which halogen is used as an antiseptic? How does it work?
- f. What is the colour change when chlorine displaces bromine?
- g. How the halogen acids are ionized in water?
- h. Why HF is weaker acid than HCl?
- i. Describe a simple chemical test that could be used to distinguish between aqueous solutions of potassium bromide and potassium iodide. Include the reagents and expected observations.
- j. Explain the chemical principles behind the use of chlorine as a disinfectant in water purification. Include relevant chemical equations in your explanation.
- k. Describe one significant disadvantage associated with the use of chlorine in water purification.
- 1. What is disproportionation reaction? Give an example.
- 2. Chlorine gas reacts differently with sodium hydroxide solution depending on the temperature and concentration.
- 3. Write balanced chemical equations for the reaction of chlorine (Cl₂) with: i) Cold, dilute sodium hydroxide (NaOH). ii) Hot, concentrated sodium hydroxide (NaOH).
- 4. For each reaction in question 3, identify the oxidation states of chlorine in the reactant (Cl₂) and in each of the chlorine-containing products. Use these oxidation states to explain why both reactions are classified as disproportionation reactions.
- a. Identify and briefly explain three major natural sources of air pollutants.
- b. How can deforestation impact air quality?
- c. Explain the reasons for the temperature trends observed in the troposphere and the stratosphere.
- d. Describe four significant anthropogenic (human-caused) activities that contribute to the deterioration of air quality. For each activity, name at least one major pollutant released.
- e. What are the environmental impacts of persistent organic pollutants (POPs)?
- f. How does polycyclic aromatic hydrocarbon (PAHs) affect human health?
- g. What is photochemical smog? Under what conditions, it forms?
- h. What type of data do air quality index (AQI) system provide?
- i. Distinguish between PM₁₀ and PM₂.₅, specifying the size ranges and describing why PM₂.₅ is generally considered more harmful to human health.
- j. What are the main chemical processes involved in the formation of acid rain?
- k. What are the specific measures to control smog?
- l. How does a catalytic converter reduce harmful vehicle emissions?
- m. Describe the sources of lead and mercury pollution.
- a. Why is there a need to crystallize a crude product?
- b. What is the function of fluted filter paper?
- c. What is the difference between a Gooch crucible and Sintered glass crucible?
- d. What type of mixtures are filtered through a Gooch crucible?
- e. What is function of fractionating column during fractional distillation?
- f. What is the stationary phase in the paper chromatography?
- g. What will happen during paper chromatography if the components of the mixture have a comparable attraction for the stationary phase?
- h. What is meant by the term developing the chromatogram in paper chromatography?
- i. What is the basic principle of paper chromatography?
- j. Why water is not generally used as a solvent in paper chromatography?
- k. Differentiate between adsorption and partition chromatography.
- l. How can you check the purity of a compound with the help of paper chromatography?
- m. You have prepared a solid sample of glucosazone in the laboratory. How will you proceed to check the purity of the sample?
- a. For which type of titration, methyl orange is used as an indicator?
- b. Explain why phenolphthalein is a suitable indicator for the titration of a weak acid with a strong base but not for the titration of a strong acid with a weak base.
- c. Explain why different indicators change color over different pH ranges.
- d. It is always advisable to use dilute solutions while performing experiments in volumetric analysis? Give a reason.
- e. White precipitates are formed when Ca²⁺, Al³⁺ and Zn²⁺ react separately with NaOH solution. How will you detect which basic radical is present?
- f. How Fe²⁺ can be distinguished from Fe³⁺ chemically?
- g. Why does Ca²⁺ not give precipitate with aqueous ammonia?
- h. How will you find out the concentration of acetic acid in vinegar solution?
- i. What precautions you need to observe while diluting a concentrated acid?
- j. Why does an aqueous solution of Na₂CO₃ behave like a base?
- k. If an aqueous solution of NaOH is kept in an open container, what changes do you expect to take place with the passage of time?
CHAPTER 1
CHAPTER 2
CHAPTER 3
CHAPTER 4
CHAPTER 5
CHAPTER 6
CHAPTER 7
CHAPTER 8
CHAPTER 9
CHAPTER 10
CHAPTER 11
CHAPTER 12
CHAPTER 13
CHAPTER 14
CHAPTER 15
CHAPTER 16
Long Questions
- Write equations for the reactions of Na and Mg with oxygen, chlorine, and water. Compare the reactivity of both elements with these in terms of metallic character.
- Explain with the help of equations, acidic and basic behavior of oxides and chlorides.
- Describe the factors affecting and periodic trends of electron affinity.
- Define ionization energy. Discuss the factors affecting and periodic trends of ionization energy.
- What are quantum numbers? Describe briefly principal and spin quantum numbers.
- Draw the shapes of s, p and d-orbitals. Justify these by keeping in view the azimuthal and magnetic quantum numbers.
- What do you mean by successive ionization energies? How the electronic shell structure of magnesium (Mg) is derived from the successive ionization energies?
- How the bonding in the following molecules can be explained with respect to valence bond theory? i) Cl₂ ii) O₂ iii) N₂ iv) HF v) H₂S
- What are the postulates of VSEPR model? Discuss the structures of the following species with reference to this theory. i) CH₄ ii) NH₃ iii) H₃O⁺ iv) PCl₅ v) SO₂ vi) SF₆
- Explain the orbital hybridization for CH₄, NH₃, BF₃, and BeCl₂.
- Draw the molecular orbital diagrams of the following molecules. Calculate their bond orders. i) H₂ ii) He₂ iii) N₂ iv) O₂
- Discuss the formation of F₂ molecule in the light of Lewis concept, VBT, and MOT.
- Differentiate limiting and non-limiting reactants. How a limiting reactant is determined from a balanced chemical equation and given data?
- Differentiate actual and theoretical yields. Why the theoretical yield is always greater than actual yield?
- A solution of sodium hydroxide (NaOH) is prepared by dissolving 2.00 g of solid sodium hydroxide in water to make a final volume of 230 cm³. a) Determine the molar mass of sodium hydroxide. b) Calculate the number of moles of sodium hydroxide used. c) Calculate the concentration of the sodium hydroxide solution in mol dm⁻³. d) If more water is added to the above solution to raise the volume of solution to 500 cm³, what would be the concentration now?
- Ammonia gas (NH₃) reacts with oxygen gas (O₂) according to the following balanced equation: 4NH₃(g) + 5O₂(g) → 4NO(g) + 6H₂O(g) In an experiment, 34.0 g of ammonia is reacted with 96 g of oxygen. a) Determine the limiting reactant. b) Calculate the maximum mass of nitrogen monoxide (NO) that can be formed. c) Calculate the mass of the excess reactant remaining after the reaction is complete. (Relative atomic masses: H = 1.0, N = 14.0, O = 16.0)
- When iron(III) oxide (Fe₂O₃) reacts with carbon monoxide (CO) in a blast furnace, iron metal (Fe) is produced according to the following equation: Fe₂O₃(s) + 3CO(g) → 2Fe(s) + 3CO₂(g) If 1.00 kg of iron(III) oxide is reacted with excess carbon monoxide, and 650 g of iron is obtained, what is the percentage yield of iron? (Atomic mass of O = 16.0, Fe = 55.8)
- 1.5g of C₂H₆ is burnt in excess of O₂ to produce CO₂ and H₂O. What volume of CO₂ is produced at STP. 2C₂H₆ + 7O₂ → 4CO₂ + 6H₂O Also calculate: a) No. of molecules of solid CO₂ produced. b) No. of O₂ molecules reacted c) No. of CH bond of C₂H₆ are broken in this reaction.
- What are London dispersion forces? Give examples, and discuss the factors affecting these forces.
- Hydrogen bonding is present in H2O, NH3, HF, (CH3)2CO and CHCl3 molecules. Sketch structures and discuss briefly.
- Discuss the structural changes when water turns into ice. Justify the empty spaces in its crystals as compared to H2O at 4°C and lower density of ice.
- How liquid crystals resemble liquids and solids? Give their uses in daily life.
- Describe the following properties of crystalline solids. i) Geometrical shape ii) Melting point iii) Cleavage plan iv) Habit of a crystal
- A sample of an unknown gas has a mass of 0.560 g. It occupies a volume of 2.87 × 10⁻⁴ m³ at a temperature of 300 K and a pressure of 1.01 × 10⁵ Pa. Calculate the molar mass of the gas. (Gas constant, R=8.31 JK⁻¹mol⁻¹)
- In a laboratory experiment, 150 cm³ of a volatile liquid was completely vaporized at 98 °C and a pressure of 1.01 × 10⁵ Pa. The mass of the vapor was found to be 0.495 g. Determine the molecular mass of the liquid (R = 8.314 JK⁻¹ mol⁻¹)
- State and explain Hess' law. Give its two applications.
- What is lattice energy? How does Born-Haber cycle help to calculate the lattice energy of NaCl?
- When 0.400 g NaOH is dissolved in 100.0 g of water, the temperature rises from 25.00 to 26.03°C. Calculate: i) q_soln (ii) ΔH for the solution process
- By applying Hess' law, calculate the enthalpy change for the formation of an aqueous solution of NH₄Cl from NH₃ gas and HCl gas. The results for the various reactions are as follows: (i) NH₃(g) + aq → NH₃(aq) ΔH = -35.16 kJ mol⁻¹ (ii) HCl(g) + aq → HCl(aq) ΔH = -72.41 kJ mol⁻¹ (iii) NH₃(aq) + HCl(aq) → NH₄Cl(aq) ΔH = -51.48 kJ mol⁻¹
- Calculate the heat of formation of ethyl alcohol from the following information: (i) Heat of combustion of ethyl alcohol is -1367 kJ mol⁻¹ (ii) Heat of formation of carbon dioxide is -393.7 kJ mol⁻¹ (iii) Heat of formation of water is -285.8 kJ mol⁻¹
- Calculate the entropy change of the surroundings ΔS°_surroundings for the reaction at 298K: 2Ca(s) + O₂(g) → 2CaO(s) ΔH°_reaction = -1270.2 kJ mol⁻¹
- For the reaction: CaSO₄(s) → Ca²⁺(aq) + SO₄²⁻(aq) Calculate ΔH°, ΔS° and ΔG° at 25°C using the following data; and discuss its spontaneity. Enthalpy of Formation: ΔH°(CaSO₄(s)) = -1432.7 kJ, ΔH°(Ca²⁺(aq)) = -543.0 kJ, ΔH°(SO₄²⁻(aq)) = -907.5 kJ Standard entropy: S°(CaSO₄(s)) = 106.7 J/K, S°(Ca²⁺(aq)) = -55.2 J/K, S°(SO₄²⁻(aq)) = +17.2 J/K
- Relate the order of a reaction to the rate law for the reaction. How do you distinguish between zero order, first order and second order reaction?
- How do you find the numerical value of a rate constant by initial and half-life methods?
- How does the activation energy profile of an uncatalyzed reaction compare with that of the catalyzed reaction?
- The reaction between hydrogen peroxide (H₂O₂) and iodide ions (I⁻) in acidic solution is believed to occur via the following mechanism: Step 1: H₂O₂ + I⁻ → H₂O + OI⁻ (slow) Step 2: OI⁻ + H⁺ → HOI (fast) Step 3: HOI + I⁻ + H⁺ → H₂O + I₂ (fast) i) Write the overall balanced equation for the reaction. ii) Identify any intermediates and catalysts in this mechanism. iii) What is the rate-determining step? iv) Write the rate equation for the reaction, expressing it in terms of the reactants in the overall reaction.
- Calculate the reaction rate if the concentration of A is 0.5 M, the concentration of B is 0.2 M and the rate constant k is 4.0 M⁻²s⁻¹. Given the rate law for a reaction: Rate = k[A][B]².
- A first order reaction is found to have a rate constant, k = 5.5 × 10⁻¹⁴ s⁻¹. Find the half-life of the reaction. Using the given temperature and rate constant data, analyze the reaction kinetics . Temp. (K) | Rate constant (cm³ mol⁻¹S⁻¹) 500 | 6.814 × 10⁻⁴ 550 | 2.64 × 10⁻² 600 | 0.56 × 10⁰ 650 | 7.31 × 10⁰ 700 | 66.67 × 10⁰
- Three experiments that have identical conditions were performed to measure the initial rate of the reaction. 2HI(g) → H₂(g) + I₂(g) Experiment | [HI] (M) | Rate (M/s) 1 | 0.015 | 1.1 × 10⁻³ 2 | 0.030 | 4.4 × 10⁻³ 3 | 0.045 | 9.9 × 10⁻³ Write the rate law for the reaction. Find the value and units of the specific rate constant, k.
- Define and explain the law of mass action and derive the expression for the equilibrium constant.
- Write the expressions for Kc for the following reactions: i) Sn²⁺(aq) + 2Fe³⁺(aq) ⇌ Sn⁴⁺(aq) + 2Fe²⁺(aq) ii) Ag⁺(aq) + Fe²⁺(aq) ⇌ Fe³⁺(aq) + Ag(s) iii) N₂(g) + O₂(g) ⇌ 2NO(g) iv) 4NH₃(g) + 5O₂(g) ⇌ 4NO(g) + 6H₂O(g) v) PCl₅(g) ⇌ PCl₃(g) + Cl₂(g)
- Write down the Kc for the following reactions. Suppose that the reaction mixture in all the case is V dm³ i) CH₃COOH + CH₃CH₂OH ⇌ CH₃COOC₂H₅ + H₂O ii) 2HI ⇌ H₂ + I₂ iii) N₂ + 3H₂ ⇌ 2NH₃
- In the equilibrium PCl₅(g) ⇌ PCl₃(g) + Cl₂(g) ΔH = 90 kJ mol⁻¹ What is the effect on the following changes? Explain your answer. i) if temperature is increased ii) volume of the container is decreased iii) catalyst is added iv) chlorine is added
- Synthesis of ammonia by Haber’s process is an exothermic reaction N₂(g) + 3H₂(g) ⇌ 2NH₃(g) ΔH = -92.46 kJ What should be the possible effect of change of temperature at equilibrium stage?
- Kc for the following reaction is 0.016 at 520°C 2HI(g) ⇌ H₂(g) + I₂(g) The equilibrium mixture contains HI = 0.08 M, H₂ = 0.01 M and I₂ = 0.01 M. To this mixture, more HI is added so that its new concentration is 0.096 M. What will be the concentrations of HI, H₂ and I₂ when equilibrium is re-established?
- The equilibrium constant for the reaction between acetic acid and ethyl alcohol is 4. A mixture of 3 moles of acetic acid and 1 mole of ethyl alcohol is allowed to come to equilibrium. Calculate the amount of ethyl acetate present at equilibrium.
- Study the equilibrium H₂O(g) + CO(g) ⇌ H₂(g) + CO₂(g) i) Write the expression of Kp. ii) When 1.00 mole of steam and 1.00 mole of CO are allowed to reach equilibrium, 33.3% of equilibrium mixture is hydrogen. Calculate the value of Kp. State the units of Kp.
- Describe the Bronsted-Lowry theory of acids and bases. Provide examples of conjugate acid-base pairs and explain clearly their relationship.
- Define the Lewis theory of acids and bases. How does this theory differ from the Bronsted-Lowry theory? Give examples of Lewis acids and bases that do not involve proton transfer.
- Discuss applications and implications of the common ion effect in various fields.
- What is the solubility product for sparingly soluble salts. Give its two applications.
- Describe the general shape of a titration curve for a strong acid titrated with a strong base. How can you identify the equivalence point on this titration curve?
- A buffer solution has a pH of 5.0. It is made from a weak acid HA with a pKa of 4.8. What is the ratio of the concentration of the conjugate base [A⁻] to the concentration of the weak acid [HA] in this buffer?
- Calculate the solubility of a sparingly salt lead (II) iodide (PbI₂) in water. It has Ksp = 1.4 × 10⁻⁸.
- The molar solubility of silver chromate (Ag₂CrO₄) in pure water at 298 K is 6.5 × 10⁻⁵ mol dm⁻³. Calculate the Ksp of silver chromate at this temperature.
- How electrode potential varies with concentration of an aqueous solution? Use the NERST equation to explain this variation.
- How Avogadro's number can be derived using an electrolytic cell?
- Describe the construction and working principle of the Zn-Cu Galvanic cell.
- What is meant by Standard Hydrogen Electrode (SHE)? How it is used to measure the electrode potential of another electrode?
- Calculate the electrode potential for a zinc electrode immersed in a 0.010 mol dm⁻³ solution of zinc sulfate (ZnSO₄) at 298 K. The standard electrode potential (E°) for Zn²⁺(aq) + 2e⁻ ⇌ Zn(s) is -0.76 V. (Gas constant, R = 8.31 J K⁻¹ mol⁻¹, Faraday constant, F = 96500 C mol⁻¹)
- A constant current of 2.00 A is passed through a solution of copper(II) sulfate (CuSO₄) for 30.0 minutes. Calculate the mass of copper deposited at the cathode. (Molar mass of Cu = 63.5 g mol⁻¹, Faraday constant, F = 96500 C mol⁻¹)
- A galvanic cell consists of a standard hydrogen electrode (SHE) and a Ni²⁺(aq)/Ni(s) half-cell. The measured cell potential at 298 K is 0.25 V, and the nickel electrode is the negative terminal. a) Write the balanced overall cell reaction. b) Determine the standard electrode potential (E°) of the Ni²⁺(aq)/Ni(s) half-cell. c) Identify which electrode is the anode and which is the cathode.
- Describe the free radical halogenation of methane using Cl₂ as an example.
- Describe the following methods for the preparation of alkenes: i) Dehydrohalogenation of alkyl halides ii) Dehydration of alcohols
- Describe the mechanism of electrophilic addition of hydrogen halides to alkenes. Discuss Markovnikov's rule in the context of hydrogen halide addition.
- Explain the following reactions of alkenes with examples: a) Halogenation b) Ozonolysis c) Epoxidation d) Polymerization
- Explain the preparation and basicity of ammonia.
- How oxides of nitrogen (NOâ‚“) cause the formation of photochemical smog and PAN? Give its mechanism.
- Give flowsheet diagram and equations involved in the contact process.
- Discuss sulfuric acid as an oxidizing agent and a dehydrating agent with three reactions for each.
- Describe and explain the relative thermal stabilities of the halogen hydrides in terms of bonds strength.
- Discuss the relative reactivity of the halogen elements as oxidizing agents. Arrange F₂, Cl₂, Br₂, I₂ in increasing order of the oxidizing power.
- Discuss the reducing power of halide ions with relevant reactions. Also explain the factors affecting it.
- Discuss sources and effects of following air pollutants on environment: i) Heavy metals ii) VOCs iii) PAHs iv) POPs
- Write short notes on the following: i) CFCs and ozone layer depletion ii) Greenhouse effect and global warming
- How the fossil fuel burning causes acid rain? Discuss in detail with chemical reactions.
- What is meant by air quality AQI? Describe the factors affecting the air quality.
- Differentiate simple distillation and fractional distillation in construction and applications.
- Describe the criterion to choose the suitable solvent in the process of crystallization.
- What is the basic principle of paper chromatography? Give its three practical life applications.
- Describe common types of the Chemistry lab hazards with two examples in each case.
- What are common accidents in the Chemistry lab? How they are managed in first aid treatment.
- How the following acid radicals are indicated and confirmed in salt analysis: i) CO₃²⁻ ii) Cl⁻ iii) NO₃⁻ iv) SO₄²⁻
- How the following basic radicals are indicated and confirmed in salt analysis: i) Cu²⁺ ii) Al³⁺ iii) Fe³⁺ iv) Zn²⁺
CHAPTER 1
CHAPTER 2
CHAPTER 3
CHAPTER 4
CHAPTER 5
CHAPTER 6
CHAPTER 7
CHAPTER 8
CHAPTER 9
CHAPTER 10
CHAPTER 11
CHAPTER 12
CHAPTER 13
CHAPTER 14
CHAPTER 15
CHAPTER 16
Administrator: SADIA ANWAR Assistant: KHADIJA AHMED
Time Allowed: ?
Total Marks: ?
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INSTRUCTONS
*Note: Write answers neatly and in sequence.Use black or blue pen only.
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